Energy changes in a reaction Notes

What Is Activation Energy, and Why Is It Necessary?

Activation Energy – the Energy Barrier

1. It acts like an energy barrier between reactants and products.
2. Only particles with enough energy (and the right orientation) can collide and react.

Why Is Activation Energy Necessary?

1. Most reactions involve:
   1. Breaking existing bonds in the reactants, and
   2. Forming new bonds in the products.
2. Breaking bonds always requires energy, so:
   1. Reactant particles must collide with enough energy to start breaking bonds.
   2. This required energy input is the activation energy.
3. Even if a reaction releases energy overall (is exothermic), it still needs an initial “push”.

Factors That Affect Activation Energy

1. Temperature
   1. Higher temperature → particles have more kinetic energy.
   2. More particles have energy ≥ $E_a$, so more successful collisions occur per second → reaction rate increases.
2. Catalysts
   1. A catalyst provides an alternative reaction pathway with a lower activation energy.
   2. More particles can now overcome the smaller barrier, so the reaction occurs faster.
   3. Importantly, a catalyst does not change $\Delta H$ (the overall energy change), only the rate.

How Do Breaking and Making Bonds Relate to the Overall Energy Change?

During a chemical reaction, two key processes occur:

1. Bonds in reactants are broken.
2. New bonds in products are formed.

Each of these steps involves energy.

Breaking Bonds – Endothermic

1. To break a chemical bond, energy must be supplied.
2. This energy is called the bond dissociation energy.
3. Bond breaking is always an endothermic process (energy is absorbed from the surroundings).

Making Bonds – Exothermic

1. When new bonds form, energy is released.
2. Bond formation is an exothermic process.
3. The surroundings often warm up when many strong bonds are formed.

Overall Energy Change ($ΔH$)

1. The overall energy change of a reaction, $ΔH$, depends on the balance between:
   1. Energy absorbed to break bonds, and
   2. Energy released when new bonds form.
2. We can think of it as: $$\Delta H=\text { Energy to break bonds }- \text { Energy released when bonds form }$$
3. If more energy is absorbed than released → $\Delta H > 0$ → endothermic reaction.
4. If more energy is released than absorbed → $\Delta H < 0$ → exothermic reaction.

Why Is This Important?

Understanding bond energies helps us:

1. Predict whether a reaction will release or absorb energy.
2. Explain why some reactions need heating while others release heat spontaneously.
3. Design processes in industry, biochemistry, and energy production.

Using Experiments to Estimate Energy Changes

We can measure and compare energy changes by monitoring temperature during a reaction.

Measuring Temperature Change

In a simple experiment:

1. Measure the initial temperature of the reaction mixture (usually a solution).
2. Allow the reaction to occur in an insulated container (e.g. polystyrene cup).
3. Measure the maximum or minimum temperature reached.
   1. If the temperature increases, the reaction is exothermic.
   2. If the temperature decreases, the reaction is endothermic.
4. Temperature vs time graphs can show how temperature changes as the reaction proceeds, then levels off.

Calculating the Energy Change (q)

The heat energy change, $q$, can be estimated using: $$q=m c \Delta T$$

where:

1. $q$ = heat energy (in joules, J)
2. $m$ = mass of the solution (in grams, g)
3. $c$ = specific heat capacity (for water ≈ 4.18 J g⁻¹ °C⁻¹)
4. $\Delta T$ = temperature change (°C)

Comparing Energy Changes Between Reactions

1. If you perform similar experiments for different reactions:
   1. Same mass of solution
   2. Same conditions
2. Then:
   1. A larger $\Delta T$ (temperature change) indicates a larger energy change ($|q|$).
   2. You can compare which reaction is more exothermic (greater rise) or more endothermic (greater fall).

Why Does This Matter?

Knowing energy changes helps us to:

1. Design heating packs or cold packs.
2. Plan energy-efficient industrial processes.
3. Understand energy transfers in biological systems and the environment.